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Atoms, Molecules and Ions
Chapter 2 Atoms, Molecules and Ions
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History of Chemistry Greeks Alchemy
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Theory vs. Law Theory – human attempts to explain or interpret natural phenomenon Law – summarizes what occurs (observed behavior)
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Dalton’s Atomic Theory
Elements - made up of atoms Same elements, same atoms. Different elements, different atoms. Chemical reactions involve bonding of atoms
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Law of Definite Composition
A compound always contains the same proportion of elements by mass
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Law of Multiple Proportions
Compounds form from specific combinations of atoms H2O vs H2O2
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Chemical Bonds Holds compounds together Need to be broken for chemical and physical changes to occur
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Periodic Table Rows (Left to Right) - periods Columns (top to bottom) - groups
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The Atom Made up of: Protons – (+) charged Electrons – (-) charged
neutrons
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Periodic Table Atomic Mass Atomic Number Alkali Metals (Group I)
Alkali Earth Metals (Group II) Chalcogens (Group VI) Halogens (Group VII) Noble Gases (Group VIII) Transition Metals
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Periodic Table Alkaline Metals – Grps. I & II Transition Metals
Non-metals Halogens – Group VII Noble Gases –Group VIII - little chemical activity
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Periodic Table Lanthanides Actinides
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Atomic Mass - # at bottom
Periodic Table Atomic Mass - # at bottom how much element weighs Atomic Number - # on top gives # protons = # electrons
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Periodic Table Atomic Mass
number below the element not whole numbers because the masses are averages of the masses of the different isotopes of the elements
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Ions Are charged species Result when elements gain electrons or lose electrons
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2 Types of Ions Anions – (-) charged Cations – (+) charged Example: F-
Example: Na+
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Highly Important! Gain of electrons makes element (-) = anion Loss of electrons makes element (+) = cation
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Charges When elements combine, they have to be in the form of IONS.
Cations and anions combine to form compounds. For a neutral compound, the sum of the charges must be ZERO. For a polyatomic ion, the sum of the charges must equal the charge of the ION.
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Examples In CO2, the charge of C is + 4 In CO, the charge of C is +2.
In KMnO4, since the charge of K is +1, O is -2 so -2 x 4 = -8, Mn must be +7. In (PO4)3-, the charge of O is -2, so -2 x 4= -8, then P must have a charge of +5, so the sum when the charges are added will be -3.
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Isotopes Are atoms of a given element that differ in the number of neutrons and consequently in atomic mass.
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Example Isotopes % Abundance 12C % 13C % 14C 11C
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For example, the mass of C = 12. 01 a. m
For example, the mass of C = a.m.u is the average of the masses of 12C, 13C and 14C.
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Determination of Aver. Mass
Ave. Mass = [(% Abund./100) (atomic mass)] [(% Abund./100) (atomic mass)]
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Take Note: If there are more than 2 isotopes, then formula has to be re-adjusted
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Sample Problem 1 Assume that element Uus is synthesized and that it has the following stable isotopes: 284Uus (283.4 a.m.u.) 34.6 % 285Uus (284.7 a.m.u.) 21.2 % 288Uus (287.8 a.m.u.) %
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Solution Ave. Mass of Uus = = 97.92 + 60.36 + 127.21
[284Uus] (283.4 a.m.u.)(0.346) [285Uus] +(284.7 a.m.u.)(0.212) [288Uus] +(287.8 a.m.u.)(0.4420) = = a.m.u (FINAL ANS.)
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Periodic Table Mendeleev – arranged elements in the (.) table
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Periodic Table Atomic Mass
number below the element not whole numbers because the masses are averages of the masses of the different isotopes of the elements
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For example, the mass of C = 12. 01 a. m
For example, the mass of C = a.m.u is the average of the masses of 12C, 13C and 14C.
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Oxidation Numbers Is the charge of the ions (elements in their ion form) Is a form of electron accounting Compounds have total charge of zero (positive charge equals negative charge)
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Oxidation States Are the partial charges of the ions. Some ions have more than one oxidation states.
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Oxidation States - generally depend upon the how the element follows the octet rule Octet Rule – rule allowing elements to follow the noble gas configuration
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Nomenclature - naming of compounds
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Periodic Table Rows (Left to Right) - periods Columns (top to bottom) - groups
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Rule 1 – IONIC COMPOUNDS Name metal or first element as is
Metals w/ Fixed Oxidation States Name metal or first element as is - Anion always ends in “–ide”
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Terminal element or anion
O - oxide P - phosphide N - nitride Se - selenide S - sulfide Cl - chloride F - fluoride I - iodide Br - bromide C - carbide
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Note Only elements that come directly from the periodic table WILL end in –IDE. POLYATOMIC IONS will be named AS IS.
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Name the following: CaO - NaCl - MgO - CaS - Na3N -
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Answers: CaO - calcium oxide NaCl - sodium chloride MgO - magnesium oxide CaS - calcium sulfide Na3N - sodium nitride
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Where do the subscripts come from?
Answer: From the oxidation states of the ions. Remember: Ions are the species that combine. Target: Compounds! (No charges!)
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Second Rule II. Ionic Compounds - Metals with no fixed oxidation states (Transition Metals) except for Ag, Zn and Al Metal(Roman #) + 1st syllable + ide Use Roman numerals after the metal to indicate oxidation state
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Name the following: Copper (I) sulfide Iron (II) oxide Tin (II) iodide
Iron (III) nitride
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Answers: Copper (I) sulfide Cu2S Iron (II) oxide FeO
Tin (II) iodide SnI2 Iron (III) nitride FeN
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What about…….? Cesium hydroxide Iron (III) acetate Lithium phosphate
Aluminum Sulfite Lead (II) sulfate Silver nitrate
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POLYATOMIC IONS Consist of more than 1 element. Have charges.
Ex. SO4 2-, SO3 2-, PO4 3-,PO3 3-
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Rule 3 – Covalent Compounds
III. For Non-metals (grps IV, V, VI VII), use prefixes. Mono – 1 Hepta - 7 Di Octa - 8 Tri – 3 Nona - 9 Tetra – 4 Deca - 10 Penta – 5 Hexa - 6
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Rule 3 – Covalent Compounds (only have Non- Metals)
Name 1st element as is. Use prefix, if necessary. Prefix + 1st element + prefix + 1st syllable of anion + ide
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Name the following compounds
CO2 - carbon dioxide N2O – dinitrogen oxide SO3 – sulfur trioxide N2O5 – dinitrogen pentoxide P2S5 – diphosphorus pentasulfide CO – carbon monoxide
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Naming Acids Use hydro + 1st syllable + “- ic acid”
I. Acids without Oxygen Use hydro + 1st syllable + “- ic acid” Example: HCl = hydrochloric acid HCN = hydrocyanic acid HBr = hydrobromic acid
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II. Acids with oxygen Polyatomic “ate” converts to “ic” + acid
Polyatomic “ite” converts to “ous” + acid - H2SO sulfurous acid H2SO4 sulfuric acid HNO3 nitric acid HNO2 nitrous acid H3PO4 phosphoric acid
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Trick! If anion ends in “ – ate”, acid ends in “ – ic” Example:
HClO4 perchlorate perchloric acid HClO3 chlorate chloric acid
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Trick! If anion ends in “ – ite”, acid ends in “ – ous” Example:
HClO2 chlorite chlorous acid HClO hypochlorite hypochlorous acid
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Name the following: HBrO4 (perbromate) HBrO3 (bromate) HBrO2 (bromite) HBrO (hypobromite)
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Law of Conservation of Mass
Fundamental laws Law of Conservation of Mass Mass is neither created or destroyed Conversion from one form to another
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Determination of Aver. Mass
Ave. Mass = [(% Abund./100) (atomic mass)] [(% Abund./100) (atomic mass)]
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Take Note: If there are more than 2 isotopes, then formula has to be re-adjusted
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Sample Problem 1 Assume that element Uus is synthesized and that it has the following stable isotopes: 284Uus (283.4 a.m.u.) 34.6 % 285Uus (284.7 a.m.u.) 21.2 % 288Uus (287.8 a.m.u.) %
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Solution Ave. Mass of Uus = = 97.92 + 60.36 + 127.21
[284Uus] (283.4 a.m.u.)(0.346) [285Uus] +(284.7 a.m.u.)(0.212) [288Uus] +(287.8 a.m.u.)(0.4420) = = a.m.u (FINAL ANS.)
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Chemical Formula Gives the combining whole number ratios of the elements in a compound C6H12O6
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Structural Formula Gives the spatial arrangement of atoms in the compound Structural formula for H2O is H – O – H
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Empirical Formula Only gives the types of elements in the compound and the ratio of the elements in the formula
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Empirical Formula Does not tell exactly how many of the elements are in the compound
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Molecular Formula Gives you the exact elemental composition of the compound Formula of the compound as it would actually exist.
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EF vs. MF Sucrose or table sugar: Molecular Formula = C6H12O6
Empirical Formula = CH2O
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Sample Problem The compound adrenaline contains % C = % H = 6.56 % O = % N = 8.28 by mass. Find the empirical formula.
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Empirical Formula EF Determination when % Masses are given
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Steps to Solve for EF Step 1: Sum up all given percentages. If total equals 100%, go to step 2. If total does not equal 100, the missing % is due to one of the component elements. Step 2: Convert Mass % to grams. Step 3: Calculate moles using mole = gram/molar mass
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Empirical Formula Step 4. To get simplest ratios, divide the moles calculated by the smallest calculated mole. You must have a ration of 1 for at least one of the element. (Follow rule for rounding). Step 5. You now have the ratios or subscripts for the EF.
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Molecular Formula Detn.
Step 1. Obtain empirical formula mass by adding atomic masses of all elements in empirical formula
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Molecular Formula Detn.
Step 2. Get ratio by applying the formula below: Molecular Formula = given molar mass Empirical formula mass
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Molecular Formula Detn.
Step 3. Multiply empirical formula subscripts by obtained ratio
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Sample Problem Caffeine, a stimulant found in coffee, contains 49.5 % C, 5.15% H, 28.9 % N, and 16.5 % O by mass. The molar mass of the compound is 195 g/mol. Determine the empirical and molecular formula of caffeine.
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Sample Problem Ibuprofen, a headache remedy, contains % C, 8.80% H, and % O by mass. The molar mass of the compound is 206 g/mol. Determine the empirical and molecular formula of ibuprofen.
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